Electronegativity

Electronegativity is one of the fundamental concepts for an understanding of chemical bonding. The first modern definition was suggested by Linus Pauling; his scale has not been improved upon since. Pauling defined electronegativity as "the ability of an atom in a molecule to attract electrons to itself."

The concept of electronegativity is especially important for a qualitative understanding of the chemical bonds—especially ionic and covalent bonds—between different types of atoms. The type of bond formed is largely determined by the difference between the electronegativities of the atoms involved. A knowledge of the electronegativities of atoms also allows us to estimate the polarity of a chemical bond and, when taken together with molecular geometry, the polarity of a molecule. Given that most chemical reactions involve the polarity of bonds in some way, electronegativity lies at the heart of chemistry. The opposite of electronegativity is termed electropositivity.

Basic concepts

The neutral atoms of different elements have differing abilities to gain or lose electrons. These properties are known as the electon affinity and ionization energy of a given element, and they can be quantitated experimentally. Electron affinity of an element is a measure of the energy released (or gained, in some cases) when one electron is added to an atom of that element. Ionization energy is the energy needed to remove an electron from an atom of that element. Atoms that attract electrons more strongly have relatively higher ionization energy and electron affinity, and they tend to form monatomic ions with a negative charge. They tend to be the atoms of nonmetals. Atoms that attract electrons more weakly have lower ionization energy and electron affinity, and they form ions with a positive charge. They tend to be the atoms of metallic elements.

Given that electronegativity is based on the degree to which an atom attracts electrons, it can be seen as related to electron affinity and ionization energy. In a covalent bond between two atoms of two different elements, the electrons in the bond will be more stable when closer to the atom with greater attraction for electrons. Consequently, the electron cloud surrounding the two atoms becomes distorted, and the bond is said to be "polarized."

As might be expected, atoms with greater electron affinity and ionization energy have stronger attraction for the bonding electrons. In the case of electronegativity, however, the atoms are considered within the context of the chemical compound they are in, not as isolated atoms. Electronegativity, therefore, is not a property of the atom itself, though we tend to treat it as such. Rather, it depends on the state of the atom in the molecule. Consequently, the electronegativity of an element cannot be measured directly—it has to be calculated as an average, on a relative scale. Several methods have been proposed for calculating electronegativity.

Pauling scale

The most common and widely used scale for electronegativities is the Pauling scale, devised by Linus Pauling in 1932. This is the scale commonly presented in general chemistry textbooks. Pauling based his scale on thermochemical data, particularly bond energies, which allowed him to calculate differences in electronegativity between atoms in a covalent bond. He assigned a value of 4.0 to fluorine, the most electronegative element, and calculated other values with respect to that. Thus the Pauling scale runs from 0 to 4, with 4 being the most electronegative. The least electronegative element is francium. Recently, the scale was revised a little—fluorine was assigned an electronegativity value of 3.98, and some minor changes were made to other reported values.

Electronegativity trends

The trends in electronegativities of the elements are shown in the table below. In general, the degree of electronegativity decreases for the elements going down each group, and it increases across each period (from left to right). This pattern follows the general trends for the values of electron affinity and ionization energy. Moving across a period, nonmetals tend to have higher electron affinities and ionization energies; and moving down a group, the values for these properties tend to decrease. The most electronegative atoms are therefore clustered in the upper, right-hand corner of the periodic table (excluding the noble gases in group 18), and the least electronegative elements are located at the bottom left of the table.

Note that the elements are shown in colors ranging from yellow to orange to red, where light yellow is used for the least electronegative element, and deep red is used for the most electronegative element.

Qualitative predictions

If we know the difference in electronegativities (ΔEN) between the atoms of two elements, we can use that value to make qualitative predictions about the nature of the chemical bond between the atoms of those elements. When the electronegativity difference between two atoms is greater than or equal to 1.7, the bond between them is usually considered ionic; for values between 1.7 and 0.4, the bond is considered polar covalent. For values below 0.4, the bond is considered nonpolar covalent.

Electronegativity and oxidation number

Oxidation and reduction reactions take place through the transfer of electrons involved in chemical bonds. If, during the course of a reaction, an element loses electrons, it is said to have been oxidized. Conversely, if an element gains electrons, it is said to have been reduced. This loss or gain may be actual or theoretical. To follow the (actual or theoretical) loss and gain of electrons by the atoms involved in a reaction, chemists assign an oxidation number (or oxidation state) to each atom in the reactants and products. The oxidation number signifies the number of charges an atom (within a molecule or ionic compound) would have if electrons were transferred completely.[1] Essentially, this means that the electrons in a chemical bond are considered as belonging to the more electronegative atom. Thus the rules for assigning oxidation numbers are based on this concept of electronegativity.

Additional scales

Two additional scales for expressing electronegativity values are based on (a) the electron affinity and ionization energy of an atom, and (b) the size and charge of an atom.

The Mulliken scale

In 1934, shortly after Pauling proposed his approach for measuring electronegativity, Robert S. Mulliken proposed a different approach. Mulliken suggested that an atom's electronegativity should be the average value of the atom's electron affinity (EAv) and ionization energy (IEv). Mulliken electronegativities, CM, may be estimated by the following equation.[2]

CM = 0.168(IEv + EAv −1.23)

In this equation, the values for electron affinity and ionization energy (reported in electron volts) must be calculated for the atom as it exists within the molecule—they are not the experimentally determined values for the neutral atom.

The Allred-Rochow scale

In 1958, A. L. Allred and E. G. Rochow proposed a separate method, based on atomic size and charge, to calculate electronegativities. They defined electronegativity as the electrostatic force exerted by the atomic nucleus on the valence electrons (outermost electrons involved in chemical bonding). When calculated using the following equation, the electronegativity values (CAR) on this scale agree well with those on the Pauling scale.

CAR = 0.744 + 0.359Zeff/r²

where Zeff is the effective nuclear charge experienced by a valence electron, and r is the distance between the electron and the atomic nucleus (covalent radius).

Credits

New World Encyclopedia writers and editors rewrote and completed the Wikipedia article in accordance with New World Encyclopediastandards. This article abides by terms of the Creative Commons CC-by-sa 3.0 License (CC-by-sa), which may be used and disseminated with proper attribution. Credit is due under the terms of this license that can reference both the New World Encyclopedia contributors and the selfless volunteer contributors of the Wikimedia Foundation. To cite this article click here for a list of acceptable citing formats.The history of earlier contributions by wikipedians is accessible to researchers here: